observed for lithium chloride in a direction such that an increase in absorptivity would be observed with dilution of the alkali halide sample. The large increase in absorbance in going from no salt to 0.35df salt repto resents the change from free EDTA% LiEIITA complex. The absorbance change in going from 0.35Jf LiCl to 5.663f LiCl cannot be caused by a shift in equilibrium in the free EDT.1LiEDT.4 ratio with ionic atrengt'h, since there is essentially no free EDTA in any of t'he solutions. This is a t once apparent by considering the LiEDT.4 equilibrium expression. Thus a change in the equilibrium constant wit,h increasing ionic st'rength cannot be responsible for this behavior. It. is possible that this increase in absorbance with increasing salt content could be caused by an increase in refractive index of the solution. From dispersion theory, it can be shown that' t,he following expression is valid if no chemical changes occur in the absorbing species and if the damping factors and oscillator strengths of absorbing species are const'ant (3).
x ahsorpt'ivity = constant where n is t,he solution refractive index a t t'he wavelength in question.
n
Thus a decrease in refractive index with decreasing salt content would
necessitate an increase in absorptivity. Using a refractive index of water (5) of 1.39 and an extrapolated value for lithium chloride of 2.3 at 222 mp and values given for the specific gravity of lithium chloride solution (a), it is possible to obtain t,he values of the molar refraction, and from that, the index of refraction of various lithium chloride solutions a t 222 mp. Such calculations show t,hat the change in refractive index is much too small to account. for the large change in absorptivity. -4possible explanation for this change in absorptivity can be attributed to chemical changes in the solution. Thus additional complexes-i.e., Li2Y-2, LisY-, etc.,-could be formed with increasing lithium concentration, with such species having significantly different absoqhvities. While no proof of this is given, it appears consist'ent with t'he present data. N o drastic effect on the results for the titration in LiCl is caused by this increase in absorptivit,y upon dilution as is shown by Table I. However, the result's are riot' as reproducible and accurate as those obt,ained in the titrations done in KC1. This is undoubt'edly caused by a dependence of the slope of the line aft'er the end point upon the exact volumes at which the absorbance values are read. Results of Table I appear to be accurate within
10%. Better results could certainly be obtained by titrating with an EDTA solution in very pure 25% LiCl, so that no dilution is obtained in the titration. Such a technique should considerably improve the preciyion obtained. The data of Figure 5 would a h 0 suggest the possibility of titrating moderate amounts of lithium in potassium chloride using this technique. LITERATURE CITED
( 1 ) Chalmers, R. A., Analyst 79, 519 (l!54). ( 2 ) Handbook of Chemistry and Physics," 44th ed., p. 2118, Chemical
Rubber Publishing Co., Cleveland, Ohio, 1962. ( 3 ) Hansen, W. X., NAA Science Center, private communication. ( 4 ) Herrera-Lancina, M.,West, T. J., ANAL.CHEM.35, 2131 (1963). ( 5 ) Landolt-Bornstein, "Numerical Values and Functions," 6th ed., Vol. 2, Part 8, p. 566, Optical Constants, Springer Publishing Co., Berlin, 1962. ( 6 ) Satelson, Samuel, Penniall, Ralph, AKAL.CHEM.27, 434 (1955). ( 7 ) Plumb, R . C., Martell, A. E., Bersworth, F. C., J . Phys. Chem. 54, 1208 (1950). ( 8 ) Strelow, F. W. E., Stork, H., ANAL. CHEM.35. 1154 (1963). ( 9 ) Sweetsei, P. k'., Bricker, C. E., Ibid., 26, 195 (1954). (10) Whitnack, G. C., Hills, &I. E., J . Electroanal. Chern. 6 , 68 (1963).
RECEIVEDfor review March 3, 1964. Accepted June 1, 1964.
Formation of Lead Chloride Fluoride at Lower pH Values for Gravimetric Purposes RAYMOND A. BOURNIQUE and LIONEL H. DAHMER' Chemistry Department, Marquetfe University, 12 I7 W. Wisconsin Ave., Milwaukee, Wis. 53233 Formation of PbClF by modification of the Hoffman and Lundell procedure to permit lower pH values was accomplished with sodium formate or sodium chloroacetate in place of sodium acetate as buffering agent. The auxiliary precipitant, chloride ion, was added in an appreciably decreased amount of HCI only rather than as HCI plus NaCI. These conditions permitted good gravimetric and volumetric recoveries of fluoride and satisfactory values of the Pb: CI ratio when the pH was maintained at 1.8 to 2.0. Precipitations were carried out in polythene beakers to avoid attack of glass by HF. At pH values below 1.8, results were definitely low while above pH 2.0 appreciably high values were obtained. Sodium carbonate in amounts as large as 2.0 grams did not interfere at pH
1786
ANALYTICAL CHEMISTRY
2.0. Hence the procedure seems suitable for gravimetric determination of fluorine in organic samples after sodium peroxide fusion. The presence of carbonate in such amounts causes slightly low results.
s
1911 when Starck (10) reported the use of lead chloride fluoride for the gravimetric determination of fluorine, a number of volumetric modifications (3, 6 , 6, 8, 12) of the procedure have been proposed. Most of these involve the titration of the chloride in the dissolved precipitate by a Volhard technique. At the high acidity involved in this titration, anions of weak acids such as carbonate generally do not interfere. In gravimetric work, however, lead salts of wch anions will often be found in the lead chloride fluoride. INCE
-1recent volumetric approach is that of Vrestal and coworkers (12) in which an excess of lead chloride is used to form the precipitate in a neutral solution, follon ed by complexometric titration of the excesb reagent. The procedures baied on titration of chloride in the precipitate have she\\ n appreciable deviations in results obtained. I n 1949, Kaufman (7') reported that insufficient control of pH may be responsible for this. I n studies of the Hoffman and Lundell (4,s)method, he concluded that the pH should be maintained between 4.6 and 4.7 for satisfactory results. Variations obtained with different volumetric PbClF methods were obPresent address: Department of Chemistry, Iowa State Univereity, Ames, Ion-a.
served by Saylor (9) and coworkers in 1951. They reported a range in the ratios of gram atoms of lead to chloride in such precipitates of 0.992 to 1.203. The range in the ratios of lead to fluoride in the same precipitates was given as 0.948 to 1.010. Further studies of the lead chloride fluoride volumetric technique would therefore appear to be warranted, with the aim of improving the process of precipitation involved. The development of a gravimetric approach with the same precipitate which would not be subject to difficulties in the presence of carbonate ion and other anions of weak acids would appear advantageous at times, in view of the somewhat tedious nature of the Volhard procedure. Tentative trials were therefore carried out on precipitation (of I'bClF in the presence and absence of carbonate ion by a simple modification of the Hoffman and Lundell method. The addition of larger amounts of nitric acid than normal caused the sodium acetate buffer system to reach lower p H values than the recommended range of 4.5 to 4.7. Promising recoveries of fluorine from sodium fluoride samples and from fluorocarbon compounds after peroxide fusion prompted further investigation as reported below. EXPERIMENTAL
Reagents. Lead acid solution was prepared for t h e precipitation of lead sulfate by dilution of 217 ml. of concentrated sulfuric acid with 1 liter of water. T o this was added lead sulfate formed by addition of sulfuric acid to a lead nitrate solution. T h e lead sulfate was first washed with 100-ml. portions of water by decantation five times a n d then transferred to the sulfuric acid solution. Just, before use, the latter was filtered through a sintered glass crucible after standing overnight or longer. Saturat,ed lead chloride fluoride solution was prepared according to t'he directions of Hillebrand and Lundell ( 4 ) . Test solutions were prepared from weighed quantities of reagent grade X a F dried a t 130" C. Apparatus. Polyet,hylene or polypropylene beakers and stirring rods of pliable metal encased in polyethj-lene were employed for formation of the P b C l F to avoid possible attack of IIF on glass. The beakers were covered with wat'ch glasses, the surfaces of which were protected from the solution by a thin sheet of polyethylene. T h e filtration of ;precipitates was carried out with sintered glass crucibles (medium porosity). Procedure. The method of Hillebrand and Lundell (.f) was used with t'he following changes: the buffering agent was sodium form,ste or commercial sodium chloroacetate (Dow Chemical Co.) in place of sodium acetate; the auxiliary precipitant was added in one form rather t,han two and in a smaller
amount of the active ion, chloride ion (1.70 ml. of 6X HCl rather than 3 ml. of 10% NaCl and 2 ml. of 6 N HC1); the washing of precipitate was carried out with a controlled volume of saturated PbClF solut.ion rather than an indefinite amount; and the p H during precipitation was appreciably lower. The fluoride test solutions were diluted to a volume of 250 ml. with water and 1.70 ml. of &V HCI were added. From a buret just enough 2 5 nitric acid was added to establish the desired p H after addition of all reagents. The required amount of acid for the desired p H in a given test' was determined by the use of a trial solution of identical composition to the test solut'ion in question during precipitation. For this purpose a Beckman Zeromatic pH meter was employed. Once the precipkation of a solution was started no observation was carried out until after filtration in order to avoid possible cont'amination by KCl from the reference calomel electrode. Solid lead nitrate was then added (usually 5.00 grams) followed by the buffering salt in solid form (3.00 grams of sodium formate or 4.40 grams of sodium chloroacetate). The solution was digested for one-half hour a t 70" C and overnight a t room temperature, and filtered. The p H of the filtrate was observed in every case and always agreed within 0.03-pH unit of the value expected from the amount of nitric acid employed in preparation of the test solut'ion. Washing of the precipitate was accomplished by the application of successive 5-ml. portions of saturat,ed PbClF solution to a total volume of 80 ml., followed by 2.5 ml. of cold water (5" C) as the concluding wash. study of the weight change of a typical precipitate in a crucible after application of various quantities of liquid indicated this to be more effective than the procedure advocated by Hillebrand and Lundell. The precipitate was then dried a t 140" C for two hours andweighed. From the amount, of PbClF rerovered, t,he fluoride content was calculated. This is reported in all tables of data as the gravimetric result. The volumetric fluoride value (and equivalent chloride cont,ent) as \vel1 as the amount of lead in the precipitate were obtained as follows: .ifter dissolving the precipitate in 150 ml. of 2N nitric acid, dilution was carried out to 250 nil. in a volumetric flask; separate aliquot,. of 100 ml. each were then taken. In one the chloride was determined by addition of excess standard silver nit'rat'e followed by back-titration with 0.1-V K C S S [modified T'olhard method of Caldwell and lloyer ( I ) ] using a 10-ml. microburet. Excess silver nitrate was controlled so as to require no more than 3 nil. of the KCSS solution. If appreciably more than this was used results were noticeably high. The direct' titration procedure of Swift ( 1 1 ) gave unreliable end points. The lead in the second aliquot was precipitated by the addition of 100 ml. of lead acid solution, followed by
evaporat,ion to dense fumes of SO,. Upon dilution with water this was repeated, 100 ml. of water were added, and the solution was heated to boiling. After digesting overnight a t room temperature, filtration was carried out in a Gooch crucible with asbestosm a t , the precipitate was washed, and ignited a t 600" C. From the weight of PbS04 thus obtained the millimoles of lead were calculated. The mean error for both chloride and lead determinations was checked with known quantities of reagent grade sodium chloride and lead nitrate. Deviations of t'he order of 1 p.p.t. were obtained. The general procedure was varied by changing one factor a t a time to ascertain the optimum conditions for accurate gravimetric fluoride results coupled with an acceptable value of the ratio of gram atoms of lead to chloride. The folloiving section summarizes the observations which were made. RESULTS
Effect of pH. I n order to obta,in various pH values for individual runs, differing amounts of 2 S nitric acid were added just before introduction of the buffering a g m t in each case. Typical results obtained are .summarized in Table I. I n the narrow range of p H from 1.90 to 2.00, good fluoride values and I'b:C1 ratios were obtained. Below this range, fluoride deviations became more negative as p H decreased, while above it deviations became more positive as pH increased. Effect of Temperature a n d Digestion Time. Various combination> of temperature and digestion time were studied for precipitate formation a t p H 2.00, other conditions remaining as for the d a t a reported in Table I . K h e n the initial precipitation or final digestion overnight (or both) were carried out a t 0" C, gravimetric and volumetric results for fluorine were appreciably high even with a n intermediate period of digestion for 30 minutes a t 70" C. Somewhat better values were found if hoth initial precipitation and digestion (for 30 minutes) were carried out a t 70" C followed by digest'ion overnight at the lower temperature. For these various conditions the positive deviat'ions in recovered amounts of fluorine ranged from 17 to 75 p.11.t. while the lead to chloride ratios were low (0.9915 to 0.9932). Much better results were obtained with initial precipitation at room temperature, digestion for 30 minutes a t 70" C, and final digestion at room temperature for one hour or mow. For periods of final digestion 5hortt.r than five hours, a definite trend ton-ard low values was noted, the deviations in fluorine values being over the range of -9.6 to - 1 . 5 p.p.t. VOL. 36,
NO. 9 , A U G U S T 1964
1787
Lower amounts did not adversely affect the lead to chloride ratio, but larger amounts caused this to decrease markedly. A similar effect was noted for changes in amount of lead nitrate added with a fixed quantity of hydrochloric acid. The optimum quantity of lead nitrate was observed to be 5.00 grams, which was the fixed amount used in the study of hydrochloric acid variation. The effect of lead nitrate variation was not as great as for chloride. The effects observed are summarized in Table 11. Effect of Carbonate Ion. X primary purpose of the investigation was the development of a gravimetric procedure applicable to samples prepared by fusion with sodium carbonate, or organic samples treated by sodium peroxide fusion. The effect of carbonate ion was therefore sought, the procedure of precipitation being that of Table I1 except that the amounts of 6N hydrochloric acid and lead nitrate were fixed a t the optimum values of 1.70 ml. and 5.00 grams, respectively. Good gravimetric and volumetric values were obtained as the amount of
The lead to chloride ratios remained quite good, falling between 0.992 and 0.995. With increase in the length of final digestion a t room temperature an improvement resulted up to a total of 12 hours. However, the results for 12 hours of final digestion were not very much better than for 5 or 6 hours. For the shorter time, deviations from +1.7 to -2.1 p.p.t. of fluorine were found. The range of deviations for gravimetric and volumetric results after 12 hours of final digestion was -0.4 to +0.6 p.p.t. The lead to chloride ratio remained virtually unchanged for five to 12 hours of digestion, being of the order 0.997 to 0.998. Effect of Chloride and Lead Concentrations. Variation in the amount of 6 S hydrochloric acid added as auxiliary precipitant with a fixed amount of precipitant, lead nitrate, caused the recovered fluorine derivations to vary from appreciably negative to pronounced positive values as the hydrochloric acid added was increased from 1.00 to 3.00 ml. The optimum result was obtained for addition of 1.70 ml. of the acid.
Table I.
Effect of pH on Precipitation of Lead Chloride Fluoride Using Sodium Formate Buffera
pH
Millimoles fluoride found Grav. Vol.
1 66
4 643
4 645
1 92
4 770 4 776 4 775 4.785 4.810 4.843
4 774 4 771 4 780 4.818 4.884 4.929
-1 83 --
1 98 2 00 2.32 2.65 3.19
4 728
4 723
Deviation (p.p.t.) Grav. Vol. -27 4 -9 6 -0 8 $0 4 +o 2 +2.3 +7.5 +14.5
Millimoles of lead found
Pb : C1 ratio
4 636 4 712 4 761 4 763 4 769 4.779 4.803 4,833
0 9980 0 9977 0 9972 0 9984 0 9978 0,9919 0 9834 0.9805
-27 0 -10 7 0 0 -0 6 +13 +9.2 $23.0 +32.5
a Conditions of precipitation were as follows: Amount of lead nitrate used, 5.00 grams; volume of 6K HC1 taken, 1.70 ml.; weight of buffer employed, 3 00 grams. Precipitation started at room temperature, solution digested at 70" C for 30 minutes and overnight a t room temperature. b Millimoles of fluoride taken, 4,774.
Table II.
Effect of Chloride and Lead Concentration on Precipitation at pH 2.00 from Sodium Formate Medium"
Amounts of precipitants added ml. 6N HC1 g. Pb(N08)s 1.00 t o 1.60 1.70 1 . 8 0 to 3.00
1.70 1.70
1.70
5 00 5 00 5 00 2 00 to 4 00 5 00 6 00 to 8 00
Ranee of deviations in F recovered (p.p.t.)b Grav. T'ol. -15 to +o 4 +2 to -36 to -0.8 +1 to
-3 +19 -3
+5
-16 to -4 -0.2 +4 0 to +26 -34 to -4 +0.4 +5 to +10
Pb:Cl variation 0 0 0 0 0 0
9975 to 0 9981 9963 to 0 9973 to 0 9979 9968 to 0
9981 9832 9981 9982
Precipitation a Conditione of preripitation: Amount of sodium formate, 3 00 grams. started at room temperature, initial digestion at 70" C for 30 minutes, final digestion overnight at room temperature. b Amount of fluoride taken, 4 774 mmoles.
1788
ANALYTICAL CHEMISTRY
sodium carbonate was increased from 0.25 to 2.00 grams. The average deviation in fluorine results found (parts per thousand) for five gravimetric trials with differing amounts of sodium carbonate in this range was -1.3 f 0.2. For the volumetric trials on the same solutions the average deviation was -1.0 f 0.4. The average lead to chloride ratio for these trials was 0.9979 f 0.0003. Upon increasing the sodium carbonate to 2.50 grams, the negative deviations increased appreciably and continued to do so as the amount of this salt rose to 5.00 grams. Over this range of addition the gravimetric deviations varied from -3.1 p.p.t. to -15 and the volumetric deviations from -2.3 to - 16. The lead to chloride ratio did not vary appreciably with this increase in carbonate but was somea hat less constant from trial to trial. The consistently negative deviations for the presence of carbonate ion may appear surprising, since one would anticipate coprecipitation of lead carbonate and high gravimetric results. The evolution of C 0 2 during the pH adjustment of solutions containing carbonate may result in the loss of H F . The loss seems insignificant until the amount of sodium carbonate present is of the order of 2.5 grams. However, variable activities of ions preqent may be partly responsible since eight grams of sodium nitrate produced negative deviations one third as large as did an equimolar amount of sodium carbonate. Precipitation in the presence of carbonate a t pH 4.5 definitely produces high results. The presence of 1-2 grams of sodium carbonate gave fluoride values by the Lundell and Hoffman procedure a t this pH which were 2.5 to 24 mg. greater than the actual amount of this ion taken, namely 0.0810 gram. Effect of Fluoride Ion Concentration. The results produced by changes in this variable, by precipitation a t p H 2.00, are given in Table 111. For a fluoride range from 2.2 to 4.8 mmoles the results were generally good. K i t h lesb than 2.2 mmoles present the fluoride recoveries were appreciably low. This observation is in agreement with those of other investigators. I n general the concentration of fluoride itself is a factor which may have a n appreciable effect. Conditions permitting satisfactory recoveries for macro amounts may not do so for micro quantities. Precision of Method. The precision of the pH-2.00 procedure was observed with a constant and near optimum amount of fluoride ion. For nine trials an average recovered value of 4.773 mmoles of fluoride was noted by both gravimetric and volumetric determinations, as compared to 4.774 mmoles taken. The average devia-
tion was 1 0 . 0 0 3 2 mmole for gravimetric and iO.0050 mmole for volumetric results. The average P b : C1 ratio was 0.9978 f 0.0007. Alternate Buffer Reagent. Since sodium formate may not operate with sufficient efficiency ,st a p H value of 2.00, sodium chloroacetate was tried in its place. However, this substance releases chloride ion by hydrolysis ( 2 ) . I t was therefore necessary to reduce the amount of 6 X HCI added from 1.70 ml. to 1.00 ml. to compensate for this. With this change, 4.40 grams of sodium chloroacetate gave results comparable to 3.00 grams of sodium formate by the procedure in question. Use of chloroacetate may (ensure closer p H control than sodium formate in view of the greater ionization constant of monochloroacetic acki as compared to formic acid. Employment of commercial material as buffering agent in this case caused no apparent difficulties in the control of chloride concentrations. Presumably the single lot of reagent which was employed was uniform in composition. Hence the reagent blank method of fixing the p H avoided trouble due to this factor. For more extensive work with the material the significance of purification would be sought. As the buffering agent was added last in the procedure with either material and formation of PbClF occurred, the p H changed from about 0.9 to 2.00 and remained virtually constant thereafter during digestion and filtration. This was carefully checked in a number of runs. Hence the buffering agent permits neutralization of acid to a definite, final p H in the process. DISCUSSION
The procedure which has been described for precipitation of PbClF a t a final p H of 2.00 a i t h either sodium formate or sodium chloroacetate as buffering (or neutrdizing) agent is expected to be of value for the gravimetric determination of fluorine in organic samples aftey sodium peroxide fusion. h gravimetric procedure based on the Lundell and Hoffman method of precipitation a t pH 4.6 can give results which are high by several parts per hundred under such conditions. The PbClF precipitates obtained in such cases were observed to effervesce upon solution in nitric acid, indicative of the presence of carbonate. I n our work we employed samples of the order of 0.2 to 0.4 gram containing from 40 to 76y0of fluorine. Since an aliquot of two fifths of each sample was (employed for precipitation, the maximum sodium carbonate content of the solutions in question could not have exceeded one gram. Therefore the low pH procedure
Table 111.
Effect of Fluoride Concentration on Precipitation of Lead Chloride Fluoride at pH 2.00 Using Sodium Formate Buffer"
Millimoles C1- found Millimoles FTaken Found0 543 0 532 1 229 1 243 2 162 2 155 2 748 2 742 3 727 3 722 4 774 4 767
-
Millimoles F- (vol.) 0 530 1 232 2 152 2 746 3 717 4 775
Deviation (p.p.t.) Grav. VOlT -20 3 -23 9 -11 3 -8 8 -3 2 -4 6 -2 2 -0 7 -1 3 -2 7 -1 5 +o 2
hIillimoles Pb+* found 1 228
Pb:Cl ratio 0 9979 0 9964
2 148
0 99x2
2 739 3 710 4 764
0 9980 0 9977
0 529
0
9976
a Conditions of precipitation were as follows: Weight of lead nitrate used, 5.00 grams; amount of buffer employed, 3 00 grams; volume of 6 5 HC1 taken, 1 70 mi.; temperature of initial precipitation, room temperature; solution digested at 70" C for 30 minutes, overnight at room temperature. Amount of fluoride taken, 4.774 millimoles.
of precipitation which we have described should be effective in reducing coprecipitat,ion of carbonate under such conditions. As we extend the procedure in this manner it will be of interest to explore more thoroughly the necessity of employing a buffer during precipitation and digestion. It may develop that satisfactory adjustment of p H before precipitation can be achieved by addition of strong acid only, the critical value required being directly checked by a p H meter with a reference electrode that contains no chloride ion. It is also possible that a more effective buffer than so far used is available. Thus dichloroacetic acid is appreciably stronger than monochloroacetic acid and should therefore be more effective a t p H values below 2. According to Drushel and Simpson ( 2 ) it is more stable t,han monochloroacetic acid and may decompose even less into an acid of smaller ionization constant, namely glycolic. However, it is possible that t'he rate of hydrolysis is sufficiently low for both substances to permit effective buffer action. Hence the detailed study of both systems may be of considerable significance, especially if initial precipitation is sought from a solution of constant p H rather t,han with addition of neutralizing agent' to reach a fixed value of pH. The studies here reported give further confirmation of the complexity of the lead chloride fluoride precipitation process. Good results for a limited range of fluoride concentration is only possible with careful control of a number of variables, among which pH appears to be important. The order of addition of reagents has been observed to have significance by us and by Winkler ( I S ) as well. Therefore, detailed studies of this effect appear to be in order as further development of the precipitation process a t pH 2.00 is sought. From the deviations given in Table
I, one might conclude that good results a t pH 2.00 are caused by compensating effects of adsorption, with low results at p H values below this being perhaps due partly to a favored coprecipitation of H F . High results above p H 2.00 possibly result from a greater tendency for PbC12 to be then adsorbed. However, further extensive studies may be required before the nature of the adsorption can be more fully understood or its relation to pH revealed. LITERATURE CITED
(1) Caldwell, J. R., Moyer, H. V., IND. ESG. CHEM.,ANAL.Eli. 7, 38 (1935). (2) Drushel, IV. A , , Simpson, G. S., J . Am. Chem. SOC.39, 2453, (1917). (3) Hawley, F., Ind. Eng. Chem. 18, 573 (1926). ( 4 ) Hillebrand, W., Lundell, G. E., Bright, H. A , , Hoffman, J. I., "Applied Inorganic Analysis," pp. 744-6. Wiley, Yew York, 1953. (5) Hoffman, F., Lundell, G. E., J . Res. ' V ~ t lRUT. . Std. 3, 589 (1929). (6) Kapfenberger, W.,Aluminum 24, 428 (1942). (7) Kaufman, S., ASAL. CHEM.21, 582 (1949). (8) Miller, J., Hunt, H., McBee, E., Ibid.. 19. 148 (19473. (9) Saylor,' J. 'H., heal, C. €I., Jr., Larkin, M. E., Tavener, A f . E., J'osburgh, W. C., ilnal. Cham. Acta 5 . 157 (1951j. (10) Starck, G., Z. Anorg. Chem. 70, 173 ( 1911). (11) Swift, E. H., Arcand, hf., Lutwack, R., Meier, I). J., ASAL. CHEM. 22, 306 (1950). (12) Vrestal, J., Havir, J., Brandstetr, J., Kotrly, S.,Chem. Listy 5 1 , 1677-9 (19573. (13) R'inkler, P. E., Rztll. SOC. Chim. Rel9. 70, 690-702 (1961). RECEIVED for review August 15, 1963.
Accepted April 29, 1064. Presented before the Analytiml llivision at the 144th
Meeting of the rlnierican Chemical Society, Los A4ngeles, California, 1063. Early work dcwribed in this paper was carried out \vith the aid of a Cott,rell Grant from the Research Corporatinn, Kew k'ork, S . Y. VOL. 36, NO. 9 , AUGUST 1964
1789
